Halogens

1
Halogens

• AS

2
Generally

• Oxidising agents
• Germicides
• Note Atoms are halogens
• Ions are halides
• Ions have 8 electrons by borrowing one, so single
  negative charge (F etc.)
• Not very soluble in water, more soluble in
  organic solvent, usually cyclohexane

3
Electronic configurations

• 2s22p5
• 3s23p5
• 4s2(3d10)4p5
• 5s2(4d10)5p5

4
Atomic radius increases down group because of

• More shielding / repulsion
• More shells
• (Even though) Bigger nucleus
• Ionic radius matches atomic radius

5
Boiling point

• (X2 molecules)
• Increases down G7
• Because of increased Van der Waals forces
• F Cl gases
• Br liquid
• I At solid

6
Reactivities (oxidising ability)

• ½X2 e ? X
• Reactivity of atoms decreases down group because
  of
• Stronger X-X bond
• Electron affinity not changing
• Hydration / Lattice energy decreasing making,
  say, NaI releases less energy than making, say,
  NaF

7
Displacement

• Higher-reactivity halogens will displace
  lower-reactivity halogens
• F gt Cl gt Br gt I gt (At)
• E.g. Cl2 2NaBr ? Br2 2NaCl
• Or Cl2 2Br- ? 2Cl- Br2
• Displaced bromine yellow/brown colour
• Displaced iodine brown and/or black precipitate
• Fluorine too dangerous

8
Electronegativity and polarisation

• Electronegativity decreases down group
• More shells so more shielding
• Radius increases so less nuclear attraction for
  shared electrons
• Bigger nucleus should attract electrons but it
  has less effect than other two
• Therefore hydrogen halides are decreasingly polar

9
Hydrogen halides (HCl etc.)

• All colourless gases, very soluble in water,
  dissociate well, ? strong acids
• E.g. HCl(aq) ? H(aq) Cl-(aq)
• Bond enthalpy decreases down group, easier to
  separate H and X-
• So ? strength of acid increases
• ? HI gt HBr gt HCl
• These acids are oxidising agents the hydrogen
  ion (H) can take an electron
• Note conc. HF not a strong acid bond
  dissociation enthalpy too high, diluted is
  stronger

10
Tests for ions

• Test solutions with acidified silver nitrate to
  make a silver halide
• (acidified to avoid formation of carbonates)
• F no change - AgF is soluble
• Cl white AgCl precipitate
• Br cream AgBr precipitate
• I yellow AgI precipitate

11

• General reaction
• AgNO3(aq) X(aq) ? AgX(s) NO3(aq)
• or Ag(aq) X(aq) ? AgX (s)
• Check with dilute and conc. ammonia solution
  (NH4OH)
• F no reaction anyway
• AgCl precipitate soluble in dilute or
  concentrated ammonia
• AgBr precipitate soluble in concentrated ammonia
  only
• AgI precipitate insoluble, even in concentrated
  ammonia

12
A redox reaction

• Cl2 H2O ? HCl HClO
• Note oxidation states
• One chlorine goes from 0 to -1 in HCl
• Other goes 0 to 1 in HClO
• This is a disproportionation one element is
  simultaneously oxidised and reduced.

13
Reduction by halide ions

• Iodide ion I- is very reducing
• happy to lose electron
• didnt really want it
• bromide less so, etc.
• Decreasing reducing power up the group
• Electrons lost particularly to strong oxidising
  agents (H2SO4, F2 etc.)

14
Homework

• Find out the reactions of H2SO4 with F- to I- and
  memorise

15
Halide ion reactions with concentrated H2SO4
MEMORISE!!

• Sulphur has an oxidation number of 6 in SO42
• F and Cl cannot reduce sulphur only white
  fumes of HF or HCl seen when NaF or NaCl added to
  conc. sulphuric acid
• F H2SO4 ? HF HSO4
• Cl H2SO4 ? HCl HSO4

16
(Br) NaBr H2SO4

• Br can reduce sulphur from 6 to 4, producing
  HBr, SO2, Br2
• Observe white fumes of HBr, invisible SO2
  (bubbles) and brown fumes of Br2 (or possibly
  liquid Br2)
• Br H2SO4 ? HBr HSO4
• Then
• 2HBr H2SO4 ? Br2 SO2 2H2O

17
(I) NaI H2SO4

• I can reduce sulphur from 6 to 4, then 0, then
  -2, producing HI, SO2, S, H2S, I2
• White fumes of HI, invisible SO2 and H2S
  (bubbles, distinctive smell), yellow sulphur, and
  purple fumes of I2
• I H2SO4 ? HI HSO4 Then
• 2HI H2SO4 ? I2 SO2 2H2O then

18
General reaction

• NaX(s) H2SO4(aq) ? NaHSO4(aq) HX(g)
• 
  fumes in moist air
• Or X H2SO4 ? HSO4 HX
• In these, the halide is acting as a base (proton
  acceptor), accepting H
• E.g.
• NaF(s) H2SO4(aq) ? NaHSO4(aq) HF(g)

19
Uses of chlorine and chlorate(I) compounds

• Under most conditions
• H2O Cl2 Ý HCl HClO
• This is a disproportionation, one chlorine atom
  is reduced, the other is oxidised.
• This makes chlorine water which can decompose
  photolytically (in sunlight)
• 2Cl2 2H2O ? O2 4HCl

20
Chlorination

• Water treatment
• Drinking water 0.7mg dm-3
• Swimming pools more concentrated
• Kills bacteria, especially E.coli from bottoms.

21
Reaction with NaOH

• Cl2 2NaOH Ý NaCl NaClO H2O
• This mixture of sodium chloride and sodium
  chlorate(I) is used as a bleach.